School of Science and Humanities

Department of Chemistry

SCYA5201

Chemistry of Main Group Elements

Unit - 1

Chemistry of Boranes

1. Boranes:

Preparation:

Diborane can also be obtained in small quantities by the reaction of iodine with sodium borohydride

in diglyme.

2NaBH4 + I2 → B2H6 + 2NaI + H2

On heating magnesium boride with HCl a mixture of volatile boranes are obtained.

2Mg3B2 + 12HCl → 6MgCl2 + B4H10 + H2

B4H10 + H2 → 2B2H6

PHYSICAL PROPERTIES

● Colourless and highly inflammable gas at room temperature.

● At high concentrations, it ignites rapidly in the presence of moist air at room temperature.

● It smells sweet.

● It has boiling point of about 180 K

● It is toxic gas

● It releases huge amount of energy when burnt in the presence of O2.

● It readily hydrolysed in the water to give hydrogen gas and boric acid.

Chemical Properties:

i) Diboranes reacts with water and alkali to give boric acid and metaborates respectively.

B2H6 + 6H2O → 2H3BO3 + 6H2

B2H6 + 2NaOH + 2H2O → 2NaBO2 + 6H2

ii) Action of air:

At room temperature pure diborane does not react with air or oxygen but in impure form it gives

B2O3 along with large amount of heat.

Structure and bonding in boranes and boron cage compounds- Closo, nido, arachno and

carboranes – Styx notation – Wade’s rule – Electron count in polyhedral boranes – Synthesis

of polyhedral boranes – Isolobal analogy between main group and transition metal fragments–

Boron halides – Phosphine- Boron heterocycles – Borazine.

B2H6 + 3O2 → B2O3 + 3H2O

ΔH = -2165 KJ mol-1

iii) Diborane reacts with methyl alcohol to give trimethyl Borate.

B2H6 + 6CH3OH → 2B(OCH3)3+ 6H2

iv) Hydroboration:

Diborane adds on to alkenes and alkynes in ether solvent at room temperature. This reaction is

called hydroboration and is highly used in synthetic organic chemistry, especially for anti

Markovnikov addition.

B2H6 + 6RCH=CHR → 2B(RCH-CH2R)3

v) Reaction with ionic hydrides

When treated with metal hydrides it forms metal borohydrides

vi) Reaction with ammonia:

When treated with excess ammonia at low temperatures diborane gives diboranediammonate. On

heating at higher temperatures it gives borazole.

Structure and Bonding in diborane:

In diborane two BH2 units are linked by two bridged hydrogens. Therefore, it has eight B-H bonds.

However, diborane has only 12 valance electrons and are not sufficient to form normal covalent

bonds. The four terminal B-H bonds are normal covalent bonds (two centre - two electron bond or

2c-2e bond).

The remaining four electrons have to be used for the bridged bonds. i.e. two three centred B-H-B

bonds utilise two electrons each. Hence, these bonds are three centre- two electron bonds (3c-2e).

The bridging hydrogen atoms are in a plane as shown in the figure 2.3. In diborne, the boron is

sp3 hybridised.

Fig.1. Bonding in diborane

Three of the four sp3 hybridised orbitals contains single electron and the fourth orbital is empty.

Two of the half filled hybridised orbitals of each boron overlap with the two hydrogens to form

four terminal 2c-2e bonds, leaving one empty and one half filled hybridised orbitals on each boron.

The Three centre - two electron bonds), B-H-B bond formation involves overlapping the half filled

hybridised orbital of one boron, the empty hybridised orbital of the other boron and the half filled

1s orbital of hydrogen.

Uses of diborane:

i) Diborane is used as a high energy fuel for propellant

ii) It is used as a reducing agent in organic chemist

iii) It is used in welding torches

2. Borazine:

● Borazine, also known as borazole.

● It is a polar inorganic compound with the chemical formula B3H6N3.

● In this cyclic compound, three BH units and three NH units alternate.

● The compound is isoelectronic and isostructural with benzene.

● Borazine has polar hexagonal structure containing 6 membered ring, in which B and N atoms are

arranged alternately.

● Because the similarity between the structures of borazine and benzene that borazine is called

inorganic benzene.

● In borazine, both Boron and Nitrogen are sp2 hybridized.

● Each N-atom has an empty p-orbital.

● B-N bond in borazine is dative bond, while arises from the sidewise overlap between the filled p-

orbitals of N-atom and empty p-orbitals of B-atom.

● Since borazine is isoelectronic with benzene, both the compounds have aromatic electron cloud.

● Due to greater difference in electronegativity values of B and N - atoms, the electron cloud in B3N3

ring of the borazine molecule is partially delocalized. While in the case of benzene ring, the electron

cloud is completely delocalized.

● MO calculations have indicated that electron drift from N to B is less than the electron drift from B to

N due to the greater electronegativity of the N-atom.

● In benzene molecules, C=C bonds are non-polar, while in case of B3H6N3, due to difference in

electronegativities between B and N atoms, the B-N bond is polar.

● It is due to the partial delocalization of electron clouds that bonding in the B3N3 ring is weakened.

● N-atoms retain its basicity and B-atoms retain its acidity.

● In borazine, B-N bond length is equal to 1.44 Ao, which is between calculated single B-N bonds (1.54

Ao).

● B=N bond length is 1.36 Ao.

● The angles are equal to 120o.

Preparation:

3B2H6 + 6NH3 —-> 3[BH2(NH3)2][BH4] —heat—> 2B3N3H6 +12H2

NH4Cl + BCl3 ——> Cl3B3N3H9 —NaBH4—> B3N3H6

NH4Cl + NaBH4 ——> B3N3H6 + H2 + NaCl

PROPERTIES

Colorless liquid

Volatile liquid

boiling point 84.5 0C

Melting point -58 0C

decomposes at -80 0C

Structure:

Chemical properties

2B3N3H6 + 6HCl ——> 2Cl3B3N3H9 —6NaBH4—> 2B3N3H12 + 3B2H6 + 6NaCl

USES

● Borazines are also starting materials for other potential ceramics such as boron carbonides

● Borazine can also be used as precursor to growing boron nitride thin films on surfaces, such as

nanomesh structure which is formed on rhodium

3. BORON HETEROCYCLES

If one or more CH group in Benzene and cyclopentadienyl are replaced by BH group then it is called

Boron heterocycles.

PREPARATION

● When stanno-hydration of diethynyl methane with dibutyl tin hydride gives 1,4-dihydro - 1,4-dibutyl

stanno benzene . It reacts with renal Boron Bromide to give 1-phenyl-1,4-dihydroxy-bromobenzene.

● When compound is converted to its anion by dipronatio with 30-butyl / lithium which on reaction

with FeCl2,which gives bis(1-phenyl borate benzene ) iron.

● Boron containing 𝞹 e- system have recently attracted much attention in research activities.

● 3 coordinate boron is isoelectronic and iso structural with positively carbocation with its vacant Pz

orbital.

● Boron is inherently electron deficient with strong e- acceptor nature allowing conjugation of organic

𝞹 - systems.

● Hence, it shows unique electronic and physical properties with lewis acid boron forms unique

complexes.

● This complex is important for turning the electronic structure of 3 - dimensional molecules.

● Boron takes trigonal planar geometry which can be used as a building block for constructing complex

molecules.

DIFFERENT TYPES OF HETEROCYCLES

4. ISOLOBAL ANALOG

Hoffmann developed the concept of isolobal analog to show the connection between an

organic or inorganic compound. According to this concept the molecular fragments of a cluster is said

to be isolobal when they possess some number of frontier molecular orbital with,

1. Similar shape, Symmetry, Radical extent

2. Approximately same energy

3. Same no of electrons are available for cluster bonding.

4. Isolobal fragments may not be isostructural and isoelectronic.

5. The relationship between isolobal fragments is shown by <r>

FIND OUT THE ISOLOBAL PAIRS FROM THE FOLLOWING

1) Mn(𝐶𝑂)5 →7+10→17

𝐶𝐻3 →4+3=7 Isolobal

2) Fe(𝐶𝑂)4 → 8 + 8 → 16

O→ 6 → 6 Isolobal

3) CO(𝐶𝑂)6 → 9 + 6 → 15

𝑅2Si→ 2 + 4 → 6 Non Isolobal

4) Mn(𝐶𝑂)5 → 7 + 10 → 17

Rs→ 1 + 6 → 7 Isolobal

5. BORON HALIDES

It is represented as BX3 where X= F, Cl, Br, I

SYNTHESIS

1.) When B2O3 is treated with halogens in presence of carbon, Boron halides are formed.

B2O3 + 3 C + 3 Cl2 → 2 BCl3 + 3 CO

2.) It is synthesized by reactions of CaF2 and boric oxide in the presence of H2SO4

3 CaF2+ B2O3+ 3 H2SO4 → 2 BF3+ 3CaSO4+ 3 H2O

PROPERTIES

● They are small in size with high charge density.

● They are covalent in nature.

● They act as non- electrolytes in liquid state.

● The boiling point is ranging from 13-3000C and it decreases when the atomic number increases.

● Boron halides are electron deficient because B atom has an incomplete octet structure.

● Boron atom accepts 6 electrons from 3 BX bonds and 2 more electrons from the donor compound.

CHEMICAL PROPERTIES:

i) BCl3 hydrolyzes readily to give hydrochloric acid and boric acid:

BCl3 + 3 H2O → B(OH)3 + 3 HCl

ii) Reduction of BCl3 to elemental boron is conduct commercially (see below). In the laboratory,

when boron trichloride can be converted to diboron tetrachloride by heating with copper metal:[5]

2 BCl3 + 2 Cu → B2Cl4 + 2 CuCl

6. Polyboranes:

● 𝐵4𝐻10

● 𝐵5𝐻11

● 𝐵6𝐻10

Wade provided the rule to correlate the no of frame work with e −with boron, cluster.

Later, Mingo extended the rule for the transition metal clusters.

General method for applying wade’s rule is as follows

Determine the total no of the valence electrons from the chemical formula.

Subtract 2 electrons for each B - H unit.

Divide the no of remaining electrons by 2 to get the no of skilled electron pairs (SEP)

A cluster with ‘n’ vertices and (n+1) SEP for bonding has closo structure.

A cluster with ‘B’ atom (n-1) vertices and (n+2) SEP for bonding has Nido structure.

A cluster with (n-2) vertices (i.e) (n-2) boron atoms and (n+3) SEP for bonding has Arachno

structure. (i.e) n-3 Boron atoms with (n+4)

STYX NUMBER OR NOTATION

The overall bonding in boranes can be represented by a 4digit number called Styx number.

s = no.of B-H-B bonds (1st digit)

t = no.of B-H-B bonds (2nd digit)

y = no.of B-B bonds (3rd digit)

x = no.of B𝐻2 group (4th digit)

EXAMPLE

1) 𝐵2𝐻6(Diborane) Styx number Styx Code Skeletal

s=2 2002 2 x 2 = 4𝑒−

t=0 0 x 2 = 0

y=0 0 x 0 = 0

x=2 2 x 2 = 4𝑒−

___

8𝑒−s

FRAMEWORK ELECTRONS :

The electrons involved in bonding in the formation of Boron cluster in boranes is known as

Framework electrons (or) electron

B-H-B 2𝑒−

B-B 2𝑒−

B-B-B 2𝑒−

-B𝐻2 2𝑒−

3) 𝐵5𝐻9(Pentaborane 9) Styx number Styx Code Skeletal

s=4(B H B) 4120 9 x 2 = 8𝑒−

t=1 (B B B) 1 x 2 = 2𝑒−

y=2(B B) 2 x 2 = 8𝑒−

x=3(-B𝐻2) 0 x 2 = 0

________

14𝑒−𝑠

4) 𝐵5𝐻11(Pentaborane 11) Styx number Styx Code

Skeletal

s=3(B H B) 3203 3 x 2 = 6𝑒−

t=2(B B B) 2 x 2 = 4𝑒−

y=0(B B) 0 x 2 = 0

x=3(-B𝐻2) 3 x 2 = 6𝑒−

________

16𝑒−

WADE’s - MNGO’S RULE :

● This theory provides electrons counting rules to predict the structure of clusters such as boranes &

carbones

● In case of transition metal cluster, for each transition metal to additional electrons are subtracted from

the total electron count

● This rule is otherwise called as PSEPT theory (Polyhedral skeletal electron pair theory )

𝐹

2=

𝑇𝑜𝑡𝑎𝑙 𝑣𝑎𝑙𝑒𝑛𝑐𝑒 𝑒𝑙𝑒𝑐𝑡𝑟𝑜𝑛𝑠 − 12 𝑥 𝑛

2

Eg: 𝑅ℎ6(𝐶𝑜)16 = 𝐹

2=

6 𝑥 9 + 16 𝑥 2 − 12 𝑥 6

2

= 86 − 72

2

= 7

n = 6

n +1 = 7 => closo

𝑅ℎ6(𝐶𝑜)16 => close (n+1) vertex

References

1.https://www.brainkart.com/article/Diborane_38581/#:~:text=Diborane%20is%20agas%20at%20roo

m,boric%20acid%20and%20metaborates%20respectively.

2. https://edubuzznotes.com/borazine-preparation-properties-structure/

School of Science and Humanities

Department of Chemistry

SCYA5201

Chemistry of Main Group Elements

Unit – 2

Chemistry of Main Group Elements

1. SILANES:

It is an inorganic compound with chemical formula, Si H4.

It is a colourless, flammable gas with a sharp, repulsive smell.

It is the silicon analogue of methane.

1.1 PREPARATION:

i) Silane is produced from magnesium silicide and hydrochloric acid:

Mg2Si + 4 HCl → 2 MgCl2 + SiH4

ii) It can also be synthesized by the reactions between silicon, hydrogen and silicon

tetrachloride:

Si + 2 H2 + 3 SiCl4 → SiH4 + SiH2Cl2

iii) A third route includes the reduction of silicon dioxide (SiO2) in presence of aluminium

and hydrogen:

3 SiO2 + 6 H2 + 4 Al → 3 SiH4 + 2 Al2O3

It has sp3hybridisation

The molecular structure of the silanes directly affects their physical and chemical characteristics.

The Si atoms in silanes are always sp3 hybridised, that is the valence electrons are said to be in

four equivalent orbitals derived from the combination of the 3s orbital and the three 3 p orbitals.

Three orbitals, which have identical energies and are arranged spatially in the form of a

tetrahedron.

It has 109.5 ° bond angle between the Si - H bonds.

Silanes, silicon halides, silicates, germenes and stannenes – Phosphorous halides, oxyacids of

phosphorous, phosphazenes – Sulphur halides, oxo acids of sulpur – Structural features and

reactivity of reactivity of S-N heterocycles. Interhalogen compounds: Structure and

properties of interhalogen compounds [ClF, ICl, ClF3, BrF3, IF3, ClF5, BrF5, IF5] poly

halides - pseudohalogens [cyanide,thiocyanate and azide] and Xenon compounds.

A silane molecule has only Si- H and Si- Si single bonds.

Si- H bonds are formed from the overlap of sp3 orbital of silicon with 1s - orbital of hydrogen.

Si- Si bonds are formed by the overlap of two sp3 orbitals on different silicon atoms.

The bond length is 746.0 pm for a Si- H bond and 233pm for Si- Si bond

The spatial arrangement of the bonds is similar to that of the four sp3 orbitals, when they are

arranged tetrahedrally with an angle of 109.5 ° between them.

Chemical properties:

i) SiH4 + 2O2 → SiO2 + 2H2O

ii) SiH4 + O2 → SiO2 + H2O

iii) 2SiH4 + O2 → 2SiOH2 + 2H2

iv) SiOH2 + O2 → SiO2 + H2O

USES

It is used as coupling agents to adhere glass Fibre to a polymer Matrix for stabilizing the

composite material.

Used to couple a Bio inter layer on a Titanium implant.

Use as water repellents.

1.2 SILICON HALIDES [Si X4] [X = Cl, Br, I, F]

SILICON TETRABROMIDE [Si Br4]

It is an inorganic compound with the formula Si Br4

It is a colourless liquid with suffering odour.

Undergoes hydrolysis readily with the liberation of HBr

It resembles closely as Si Cl4 in all its properties.

Synthesis:

It is synthesized by the reaction of Si with HBr at 6000C.

Si + 4HBr → SiBr4 + 2H2

Reactivity:

Reduction:

SiBr4 can be readily reduced by hydrides or complex hydrides.

4R2AlH + SiBr4 → SiH4 + 4R2AlBr

USES

The pyrolysis of SiBr4 has the advantage of deposition silicon and faster rate.

Pyrolysis of SiBr4 for followed by treatment with NH3 yeild Silicon nitride coating a hard

compound used for ceramic, scalants and production of many cutting tools.

SiF4: silicon tetrafluoride:

It is mainly occur in the volcanic fumes containing significant amount which on hydrolysis forms

of hexa fluoro silicic acid

PREPARATION

It is prepared by heating BaSiF6, above 300° C

BaSiF6 → Si F4 + BaF2

Uses

It is used in microelectronics and organic synthesis.

SiI4: Silicon tetraiodide

SiH4 + I2 → SiHI3

SiH2I2

SiHI3 + HI

SiI4

Physical properties:

Colourless liquid at room temperature

1.3 SILICATES

They are the largest most interesting and complicated class of minerals than any other minerals.

Approximately 30% of all minerals are silicates and some geologist estimate that 90% of the

earth crust is made up of silicate SiO44- base minerals.

Silicate is based on the basic chemical unit SiO44- tetrahedron shaped and anionic group.

The central Silicon Ion has a charge of positive four which each oxygen has a charge of negative

two (-2) and thus Si-O bond is equal to one half the total bond energy of oxygen.

This condition level the oxygen with the the option of bonding to another silicone ion and

therefore linking one SiO44- tetrahedron to another.

CLASSIFICATION

Orthosilicates

Pyrosilicates

Ring and chain silicates

Double chain silicates

Silicates with sheet structure

Silicates with 3D fragments

STRUCTURE AND BONDING

The main structural unit of silicate is Td cluster connecting one silicon and four O atom

The size of SiO44- tetrahedral is relative stable with Si-O Bond length verifying from 0.161 to

0.164nm.

The strength of Si-O bond is pretty high that provides the thermal stability and chemical

resistance of the majority of silicate compound.

The TetrahedralCluster can be polymerized

(i.e) linked to each other through the bridging oxygen atoms.

They are able to form polymers by means of linkage with one, two, three or four neighbouring

tetrahedra, forming siloxane, Si - O - Si bonds.

Other ions can be located in the silicate lattices such as Li+, Na +, K+, Be2+, F-etc.

Some cations such as Al, B, Be, are able to isomorphically substitute Si atoms in the Si-O

tetrahedral.

However, most of them are located out of the anionic framework and play the role of “charge

balanced cations” and they are all six- coordinates.

Uses

It is used as thermal insulators.

Used to make oscillators, used in watches, radios and pressure gages.

Used to make glass and ceramics.

Used to form microchips for watches

1.4 SULPHUR - NITROGEN COMPOUNDS

PREPARATION

PROPERTIES

Solid , melting point -1780 C

Thermochromic it changes colour with temperature

At liquid N2 temp - colourless

At room temp.orange - yellow

At 1000C- red

Crystals are reasonably stable to attack by air, but explosively sensitive to shock and friction.

STRUCTURE: Cradle Shaped Structure

1. Average S- N bond length 1.62 Å

2. S-N bond will have partially double bond character

3. All the bonds lengths are equal.

4. Therefore, there is delocalisation.

5. S-S distance is 2.58 Å (weak bonding).

Its structure is not accounted for by simple classical bonding diagram, several canonical

structures are possible.

CHEMICAL PROPERTIES

PROPERTIES:

Colourless and Crystalline in nature.

It has square structure.

Super conductors at 0.26K.

Delocalization of 𝝅 electrons takes place .

Conduction band is formed by overlapping of 𝝅Aorbital .

Unpaired electrons on sulphur atom will have half fill bond. Also called as one- dimensional

metal.

NSF3-Thiazyltrifluoride

1.5 GERMENES AND STANNENES

Ge - 32 (p- block)

Electronic configuration : (Ar) 3 d10 4 s2 4 p2

SYNTHESIS

Fluorodimesitylvinyl germene was prepared in one step from vinyl magnesium bromide and

difluoro dimesityl germane.

Subsequent addition to compound 1 of a molar equivalent of tert- butyl lithium at -780 C gave the

- lithiorgermane (2), the elimination of lithium fluoride occurred at -500 C with nearly

quantitative formation of dimesityl tertiary butyl germene(3)

PROPERTIES

Good thermal stability

Highly reactive, particularly towards electrophilic or nucleophilic reagents

REACTIVITY

1.6 Xenon Compounds:

ELEMENT OUTER

CONFIGURATION

FIRST

IONISATION

ENTHALPY

(K J mol-1)

NORMAL

BOILING

POINT

(K)

VOLUME IN

PERCENTAGE IN

ATOM

(x 104)

He2 1 s2 2369 4.2 5.2

Ne10 2 s22 p6 2078 27.1 18.2

Ar18 3 s2 3 p6 1519 87.3 9340.0

Kr36 4 s2 4 p6 1349 120.3 11.4

Xe54 5 s2 5 p6 1169 166.1 0.08

Rn86 6 s2 6 p6 1036 208.2 -

STRUCTURE

3 lone pair of electrons are present

Spectroscopy analysis has shown that, XeF2molecule has linear geometry.

This is explained by VSEPR theory

sp3d hybridisation.

2) Xe F4 - Xenon Tetraflouride:

STRUCTURE

2 lone pair of electrons are present

Spectroscopy analysis has shown that, XeF4 molecule has square planar geometry.

This is explained by VSEPR theory

sp3d2 hybridisation.

3- XeOF4 : Xenon oxytetrafluoride

STRUCTURE

it has one lone pair of electrons

it has square planar structure

SP2d2Hybridization

this is explained by VSEPR theory

1.7 PHOSPHORUS HALIDES

A phosphorus halide is a compound that a phosphorus forms with a halogen. A phosphorus

halide is of two types. They are PX3 and PX5. Here, we refer X to a halogen. It could be anything

from fluorine, chlorine, bromine or iodine. However, the most common phosphorus halide is that

of chloride. These chlorides are usually covalent in their nature.

1.7.1 Phosphorus Trichloride PCl3:

This is an oily and sleek fluid.

It is very lethal in nature.

The shape of this compound is that of a triangular pyramid.

The atom of phosphorus exhibits sp3hybridization.

As we can see in the above diagram, phosphorus has its sp3orbitals. It has only one electron and

it gives that electron to a p orbital electron from 3 chlorine atoms. The fourth sp3

Orbital is full. It is a solitary lone pair. Thus, it cannot form a bond. However, it repels alternate

bonds. This creates a state of the shape of trigonal pyramidal.

Preparation:

We obtain phosphorus trichloride by passing dry chlorine over heated white phosphorus. The

reaction that takes place is as follows:

P4+ 6Cl2 → 4PCl3

We can also obtain this compound by the reaction of thionyl chloride with white phosphorus.

Below is the reaction to it.

P + 8SOCl2 4SO2+ 2S2Cl2

Chemical Properties:

Phosphorus trichloride hydrolysis when we dampen it.

PCl3+ 3H2O →H3PO3+ 3HCL

It reacts with natural compounds having a –OH group and gives their ‘chloro’ subsidiaries as

products.

Structure of PCl3:

The phosphorus particle in the centre of PCl3 exhibits sp3 hybridisation. It has 3 bond sets and 1

lone pair of electrons. Due to this reason, it has a pyramidal shape. It acts as a Lewis base

because it has the capacity to donate its lone pair of electrons to other electron-lacking particles

or atoms.

1.7.2 Phosphorus Pentachloride PCl5:

It is yellowish-white in colour. Phosphorus Pentachloride is a very water delicate solid. It

dissolves in organic solvents like carbon tetrachloride, benzene, carbon disulphide, diethyl ether.

Its structure is that of a trigonal bi-pyramid. We find this structure primarily in vaporous and

fluid stages. In the solid state, we can find it as an ionic solid, [PCl4]+[PCl6]-. Here, the cation,

[PCl4]+ is tetrahedral and the anion, [PCl6]- is octahedral. We must know that the molecule has

three tropical P-Cl bonds and two pivotal P-Cl bonds. Due to the more prominent repulsion at

hub positions as compared to the central positions, we see that the two axial bonds are longer

than tropical bonds.

Preparation:

i)We can produce pentachloride by the reaction with an excess of dry chlorine.

P4+ 10Cl3→ 4PCl5

ii) We can also produce it by the reaction of SO2Cl2 and phosphorus.

P4 + 10SO2Cl2 → 4PCl5+ 10SO2

Structure of PCl5 As we discussed earlier, the central phosphorus atom in phosphorus pentachloride experiences

SP3d hybridisation. All the five electrons combine in these hybrid orbitals as bond sets. The

molecular shape of the particle is trigonal bipyramidal.

After the five electrons get hybridised, we get five electrons of equivalent size and shape. Three

of them frame a triangle (120° partition) in the centre. One bond is above and one is underneath

those three.

1.8 Oxoacids of Phosphorus

Few Popular Oxoacids of Phosphorus

In this section, we will look at some of the most important and popular oxoacids of phosphorus.

1.9 Sulfur Halides:

1.9.1 Sulfur hexafluoride:

Sulfur hexafluoride (SF6) is a gas at standard temperature and pressure (25 °C, 1 atm). The most

common synthesis involves the direct reaction of sulfur with fluorine yields SF6.

It should be noted that while SF6 is highly stable, SCl6 is not formed. The explanation of this

difference may be explained by a consideration of the Born-Haber cycle shown in Figure. A

similar cycle may be calculated for SCl6; however, a combination of a higher dissociation energy

for Cl2 and a lower S-Cl bond energy (Table) provide the rational for why SCl6 is not formed.

Born-Haber cycle for the formation (ΔHf) of SF6: where D(X-Y) = dissociation energy for X-Y

bond, E(S-F) = S-F bond energy, and S* indicates 6 coordinate sulfur.

Comparison of diatomic bond dissociation and S-X bond energy for the fluorine and chlorine analogs.

Bond dissociation energy kJ/mol Bond energy

D(F-F) 158 E(S-F)

D(Cl-Cl) 262 E(S-Cl)

The S-F bond length (1.56 Å) is very short and consistent with π-bonding in addition to σ-

bonding. Like SiF62-, SF6 is an example of a hypervalent molecule.

Molecular orbital bonding in SF6.

Sulfur hexafluoride is an unreactive, non-toxic compound. Its inert nature provides one of its

applications, as a spark suppressor. The hexafluoride is generally resistant to chemical attack,

e.g., no reaction is observed with potassium hydroxide (KOH) at 500 °C. The low reactivity is

due to SF6 being kinetically inert due to:

Coordination saturation precluding associative reactions with nucleophiles.

Strong S-F bond (360 kJ/mol) limiting dissociative reactions.

Thermodynamically SF6 should react with water (ΔH = -460 kJ/mol), but the rate factors are too

great. Sulfur hexafluoride can be reduced with sodium in liquid ammonia, Equation, or with

LiAlH4. In each of these reactions the mechanism involves the formation of a radical, Equation.

The reaction with sulfur trioxide yields SO2F2, Equation, however, the reactions with carbon or

CS2 only occur at elevated temperatures (500 °C) and pressure (4000 atm).

2. INTERHALOGEN COMPOUND

An interhalogen compound is a molecule which contains two or more different halogen

atoms

(fluorine, chlorine, bromine, iodine, or astatine) and no atoms of elements from any other group.

Most interhalogen compounds known are binary (composed of only two distinct elements)

The common interhalogen compounds include Chlorine monofluoride, bromine

trifluoride, iodine pentafluoride, iodine

heptafluoride, etc

Interhalogen compounds into four types, depending on the number of atoms in the particle. They

are as follows:

XY

XY3

XY5

XY7

X is the bigger (or) less electronegative halogen.

Y represents the smaller (or) more electronegative halogen.

All are of the type XX'n where n = odd nos. X' is

always the lighter halogen

since the smaller halogen X' are bonded around the larger X. As the

ratio of the radius of X to that of X' increases, n also increases

Properties of Interhalogen Compounds

A lot of these compounds are unstable solids or fluids at 298K. A few other compounds

are gases as well. As an example, chlorine monofluoride is a gas. On the other hand,

bromine trifluoride and iodine trifluoride are solid and liquid respectively.

These compounds are covalent in nature.

These interhalogen compounds are diamagnetic in nature. This is because they have bond

pairs and lone pairs

Interhalogen compounds are very reactive. One exception to this is fluorine. This is

because the A-X bond in interhalogens is much weaker than the X-X bond in halogens,

except for the F-F bond.

2.1 Pseudohalogens:

Compounds that resemble the halogen elements, X2, in their chemistry.

E.g. (CN)2 cyanogen, (SCN)2 thiocyanogen, ICN iodine cyanide.

Certain ions that have sufficient resemblance to halide ions are sometimes referred to as

pseudohalide ions.

E.g. N3-, SCN-, CN-.

It is formed by the combination of two or more P block elements with a negative charge CN-

cyanide group a combination of carbon and nitrogen with uni-negative charge. They are called

pseudohalogens as they form covalent compounds, complexes similar to the halogens. They differ

from halogens as they are able to polymerize unlike halogens and their complexes are not

paramagnetic.

The anions CN–, CNO–, N3–, OCN–, and SCN– can be coined classical linear pseudohalides. A

small species can be classified as a classical pseudohalogen when it fulfills the following criteria

with respect to a halogen-like chemical behaviour: A pseudohalogen (X) forms

a strongly bound (linear) univalent radical (X×),

a singly charged anion (X–),

a pseudohalogenhydrogen acid of the type HX,

salts of the type M(X)n with silver, lead and mercuric salts of low solubility,

a neutral dipseudohalogen compound (X–X) which disproportionates in water and can be

added to double bonds, and

interpseudohalogen species (X–Y).

While many linear pseudohalogens (e.g. CN, OCN, CNO, N3, SCN) are known, often the

corresponding pseudohalide acids, dipseudohalogens, and interpseudohalogens are

thermodynamically highly unstable (e.g. HN3, OCN–NCO, NC–SCN) with respect to N2/CO

elimination or polymerisation or indeed, remain unknown (e.g. N3–N3).

PROPERTIES OF PSEUDOHALOGENS

These are volatile.

Pseudohalogens, like the halogens, add at ethylenic double bond linkage.

Pseudohalogens , like the halogens, react with alkalies.

Pseudohalogens form covalent pseudohalides.

Eg : Thiocyanogen (SCN)2

PREPARATION:

It is prepared by Pb(SCN)2 with Br2 in the real solution at 0 ◦C.

Pb(SCN)2 + Br2 → (SCN)2 + Pb Br2

PROPERTIES:

It is yellow solid

Insoluble in water.

It oxidises Cu(I) salts to Cu(II) salts.

Sulphur is precipitated from H2S.

USES OF PSEUDOHALOGEN COMPOUNDS

• Used as an oxidising agent.

• Used for the determination of unsaturation in organic compound.

https://www.schulz.chemie.uni-rostock.de/forschung/halogenpseudohalogen-chemistry/

School of Science and Humanities

Department of Chemistry

SCYA5201

Chemistry of Main Group Elements

UNIT 3

ORGANO METALLIC REAGENTS AND MACROCYCLES

1. Organolithium Reagents:

Organolithium reagents are organometallic compounds that contain carbon – lithium bonds. These

reagents are important in organic synthesis, and are frequently used to transfer the organic group or

the lithium atom to the substrates in synthetic steps, through nucleophilic addition or simple

deprotonation.

Synthesis:

The best general method for RLi synthesis involves the reaction of an alkyl or aryl chloride with

lithium metal in benzene or an aliphatic hydrocarbon (e.g., hexane)

Metal-hydrogen exchange reactions

Properties:

Alkyl lithium compounds are either low melting solids or liquids, with high volatility

(depending on the substituent) due to the covalent nature of the bonding.

They are soluble in aliphatics, aromatics, and ethers.

Organolithium compounds react rapidly with air and water (both vapor and liquid).

Synthesis and reactivity of organo-lithium, beryllium and magnesium compounds.

Macrocyclic Ligands: Calixarines, cryptands and crown ethers-structure and applications in

complexation chemistry.

Structure:

The structure of organolithium compounds is dominated by their highly oligomeric nature as a result

of 3-center 2-electron bridging bonds. In all cases the extent of oligomerization is dependant on the

identity of the alkyl (or aryl) group. The alkyl-bridged bond is similar to those found for beryllium

and aluminum compounds. In the vapor phase any particular organolithium derivative show a range

of oligomeric structures. For example, the mass spectrum of EtLi shows ions associated with both

tetramers (e.g., [Et3Li4]+) and hexamers (e.g., [Et5Li6]

+). The structures of the different oligomers have

been predicted by molecular orbital calculations

Chemical Properties:

Hydrolysis:

Organolithium compounds react with water to give the hydrocarbon and lithium hydroxide. Lithium

alkyls also react with other hydroxylic compounds such as alcohols and carboxylic acids.

Reaction with carbonyls:

Organolithium compounds react with organic carbonyls (aldehydes, ketones, and esters) to yield the

alcohol on hydrolysis. This synthetic route is particularly useful since lithium reagents are far more

reactive than the analogous Grignard, allowing reactions to be carried out at lower temperatures and

minimizing enolization side reactions.

The high reactivity of alkyl lithium compounds means that they react with carboxylic acids to yield

the ketone rather than the lithium carboxylate.

Organolithium compounds generally react with α,β-unsaturated ketones to give the 1,2-addition

product. However, lithium dialkylcuprates, which are formed from the alkyl lithium and copper(I)

iodide, add exclusively by the 1,4-addition.

Transmetallation:

One of the most useful methods of preparing organometallic compounds is the exchange reaction of

one organometallic compound with a salt of a different metal, Equation. This is an equilibrium

process, whose equilibrium constant is defined by the reduction potential of both metals. In general

the reaction will proceed so that the more electropositive metal will form the more ionic salt (usually

chloride).

Lithium reagents may be used to prepare a wide range of organometallic compounds.

2. Grignard Reagent: RMgX

Organomagnesium compounds are less reactive than the corresponding organolithium ones because,

due to the higher electronegativity of magnesium compared to lithium, the carbon–magnesium bond

is less polarized than the carbon–lithium one.

Preparation:

R3C−X + Mg → R3C−MgX Halide reactivity in these reactions increases in the order: Cl < Br < I and Fluorides are usually not

used. The alkyl magnesium halides described in the second reaction are called Grignard Reagents after

the French chemist, Victor Grignard, who discovered them and received the Nobel prize in 1912 for

this work.

Reaction of Organometallic Reagents with Various Carbonyls:

Because organometallic reagents react as their corresponding carbanion, they are excellent

nucleophiles. The basic reaction involves the nucleophilic attack of the carbanionic carbon in the

organometallic reagent with the electrophilic carbon in the carbonyl to form alcohols.

Grignard and Organolithium Reagents will perform these reactions:

Addition to formaldehyde gives 1° alcohols

Addition to aldehydes gives 2° alcohols

Addition to ketones gives 3° alcohols

Addition to carbon dioxide (CO2) forms a carboxylic acid

Mechanism for the Addition to Carbonyls

The mechanism for a Grignard agent is shown; the mechanism for an organometallic reagent is

the same.

1) Nucleophilic attack

2) Protonation

Organometallic Reagents as Bases

These reagents are very strong bases (pKa's of saturated hydrocarbons range from 42 to 50).

Although not usually done with Grignard reagents, organolithium reagents can be used as strong

bases. Both Grignard reagents and organolithium reagents react with water to form the

corresponding hydrocarbon. This is why so much care is needed to insure dry glassware and

solvents when working with organometallic reagents.

In fact, the reactivity of Grignard reagents and organolithium reagents can be exploited to create a

new method for the conversion of halogens to the corresponding hydrocarbon (illustrated below).

The halogen is converted to an organometallic reagent and then subsequently reacted with water

to from an alkane.

Grignard reagents react with oxygen, sulfur, and halogens to form substances

containing C−OC−O, C−SC−S, and C−XC−X bonds, respectively:

3. Organoberyllium Compounds:

Organoberyllium compounds are more difficult to prepare and are less nucleophilic than their

magnesium counterparts. These features, together with their high toxicity, have limited their utility.

Organoberyllium compounds are best prepared via transmetallation reactions or by reaction of

beryllium halides with other organometallic compounds.

HgMe2 + Be → Me2Be + Hg at 383K

2 PhLi + BeCl2 → Ph2Be + 2 LiCl

In the vapour phase Me2Be is monomeric but in the solid state it is polymeric and the bonding is

considered electron deficient with 3 centre - 2 electron bonds. With higher alkyls the amount of

polymerisation decreases and the tert-butyl derivative is monomeric and linear in both solid and

vapour phases.

2NaCp+BeCl2⟶Cp2Be+2NaCl

Beryllocene:

The reaction of NaCp with beryllium chloride leads to beryllocene (Cp2Be)

2K[C5Me5]+BeCl2→(C5Me5)2Be+2KCl

reaction is carried out in an ether/toluene mix at 388K.

The solid state structure suggests that the two rings are bound to the Be differently such that 1 is

designated η5 and the other η1.

Structure of (η1−Cp)(η5−Cp)Be(η1−Cp)(η5−Cp)Be also called Beryllocene

The experimental 1H NMR spectrum adds to the confusion of the bonding since even at 163 K the

protons all appear equivalent. This is accounted for by fluxional processes. Some variations of the

compound have been prepared to see how general this effect is, for example, 4 protons on each ring

replaced by methyl groups and all 5 protons replaced by methyl groups, (meCp)2Be, and even 4 on

one ring and 5 on the other replaced with methyl groups. In the first case the fluxional process was

observed down to 183K and in the second case the two rings were found to be coparallel and staggered.

(Note that the structure of ferrocene is described as eclipsed when prepared at very low temperatures

or in the gas phase but when formed at higher temperatures it is disordered and more staggered and

since the barrier to rotation of the two rings is quite low, at 298 K in the solid state there is motion).

However, Cp*2Be possesses a sandwich structure with both the rings are coplanar.

4. Macrocyclic ligands:

A macrocyclic ligand is a macrocycle with a ring size of atleast nine or three donor sites

Eg: crown ethers

Porphyrins

Cryptands

Calixarenes, etc.

What is the need to study these molecules

• They are exceptionally versatile.

• selectively binding a range of metal ions and also a variety of organic neutral and ionic

species.

• Crown ethers are currently being studied and used in many applications beyond their

traditional place in chemistry.

4.1 Crown Ethers

• It is a type of cyclic ether consisting of a ring of carbon and oxygen atoms, with two or

more carbon atoms between each oxygen atom.

• Crown ethers are cyclic chemical compounds that consist of a ring containing several ether

groups.

• Crown – resemblance between the structure of a crown ether bound to a cation and a crown

sitting on a person’s head.

Nomenclature

• The crown ethers are designated as per IUPAC nomenclature and also by Pederson’s crown

nomenclature.

• It is always represented by notations.

• They are named as M-crown-N,

where M = total no. of atoms in the in cycle,

N = total no. of oxygen atoms.

Eg:

• Its inner side is hydrophilic (reacts with water)

• Outer part is hydrophobic in nature (repelled by water)

Thiacrowns:

It’s a special type of crown ether where the sulphur atoms takes the position of oxygen.

Synthesis of Crown ethers:

• Pederson synthesized the crown ethers by ionization of the phenolic hydroxyls.

Applications

• Synthetic application

• Analytical application

Synthetic application

• Esterification

• Saponification

• Anhydride Formation

• Potassium Permanganate Formation

• Aromatic substitution reaction

• Elimination Reactions

• Displacement reactions

• Generation of carbenes

• Formation of superoxide anion

• o,c,n-alkylations etc.

Esterification

Analytical Application:

• Determination of gold in geological samples.

• Supercritical fluid extraction of trace metal from solid and liquid materials.

• Oxidation and determination of aldehydes.

• Used in the laboratory as phase transfer catalyst.

4.2 Cryptands:

• These materials were first reported in 1969.

• Cryptands are a family of bicyclic and polycyclic multidentate ligands for a variety of

cations.

• Cryptand means this ligand binds substrate in a crypt interring the guest in a burial.

• The most common and most important cryptand is N[CH2CH2OCH2CH2OCH2CH2]3N.

The formal IUPAC name for this compound is 1,10-diaza-4,7,13,16,21,24-

hexaoxabicyclo[8.8.8]hexacosane.

• This compound is termed as [2,2,2] cryptand.

• The number indicates the number of ether oxygen atoms (binding sites) in each of the three

bridges between the amino nitrogen “caps”.

Synthesis:

• The cryptand can be synthesized by the following method for [2,2,2] cryptand.

Applications

• Used as phase transfer catalyst.

Eg when KF is solubilized in the presence of 18-crown-6 or [2,2,2]cryptand, K+ ion is stabilized

as [K(macrocycle)]+ and the unsolvated F- ion acts as a poweful nucleophile.

• They are able to bind insoluble salts into organic solvents.

Eg when KMnO4 solubilized in benzene in presence of crown-6 or suitable cryptand produces

a purple solution called purple benzene.

• Used in crystallization process.

• Used for selecting the ion selective electrodes.

• Used for seperation of chemically similar metals ions based on size factor (eg., lanthanides,

alkali metal ions etc.,)

4.3 Calixarenes:

• It is a macrocycle or cyclic oligomer obtained from hydroxyalkylation product of phenol

and aldehyde.

• Calixarenene- calix or chalice which means it resembles like a vase

arene means aromatic building block.

• It have hydrophobic cavity that can hold smaller molecules or ions like cryptands.

• Its nomenclature states that by counting the no. of repeating units in the ring.

Eg. Calix[4]arene has 4 units in the ring.

Calix[6]arene has 6 units in the ring.

Calix[4]arene with tert-butyl substitution.

Synthesis:

Applications:

• Applied in enzyme mimetics.

• Used as ion selective electrodes or sensors.

• Used a selective membranes in ourifying technology.

• Used in non-linear optics.

• Used as stationary phase in HPLC.

• Used as scaffolds in cancer immunotheraphy.

References:

1. Text book of organic synthesis special techniques-vk ahulwalia, renu Aggarwal.

2. Analytical applications of crown ethers, G.S. Vasilikiotis, I.N. papadoyannis, and Th. A.

Kouimtzis

3. Synthesis of Thiacrown and Azacrown ethers based on the spiroacetal framework by,

Marica Nikac University of Western Sydney.

4. https://chem.libretexts.org/Bookshelves/Organic_Chemistry/Map%3A_Organic_Chemist

ry_(Smith)/Chapter_20%3A_Introduction_to_Carbonyl_Chemistry_Organometallic_Rea

gents_Oxidation_and_Reduction/20.09_Organometallic_Reagents.

5. https://cnx.org/contents/9G6Gee4A@25.3:gjp7zkI5@2/Organolithium-Compounds

6. https://chem.libretexts.org/Bookshelves/Organic_Chemistry/Map%3A_Organic_Chemistry_(

Bruice)/11%3A_Organometallic_Compounds/11.01%3A_Organolithium_and_Organomagne

sium_Compounds

School of Science and Humanities

Department of Chemistry

SCYA5201

Chemistry of Main Group Elements

UNIT 4

ORGANYLS

1. Introduction:

Organoaluminium compounds are known for more than 150 years. Following the groundbreaking

experiments of Frankland (1849) on the preparation of dialkylzinc reagents by the reaction of zinc

with alkyl iodides

1.1 Preparation methods of Organoaluminium Compounds:

1.1.1 Transmetallation:

The most widely employed synthesis of organoaluminium compounds on laboratory scale is the

transmetallation reaction (salt metathesis), using organometallics and an aluminium salt. The driving

force of this reaction is the formation of a thermodynamically more stable carbon-metal bond, which

is usually the more covalent bond.

The reaction of organomagnesium reagents with AlCl3 has been investigated thoroughly as well.

Since early attempts41 this procedure has developed to a powerful synthetic methodology. Thus, for

example starting from dimesitylmagnesium, sterically hindered organoaluminiums can be prepared

easily

Preparation and reactivity of aluminium organyls – Carbalumination, hydroalumination – Chemistry of

Ga(I) and In(I) – Reduction of Al, Ga and In organyls – Germanium, tin and lead organyls.

Even functionalized organoaluminiums have been prepared recently by the reaction of

functionalized organomagnesium reagents with aluminium salts under Barbier conditions

1.1.2 Metal-Exchange Reaction:

The preparation of organoaluminium compounds has been achieved starting from mercury organyls

(Scheme 6).49 Of course, because of the use of toxic mercury this approach is not applicable

nowadays anymore. Such an exchange is not only occurring when the metal is used in its elemental

state, but also between two organometallic compounds. Thus, several groups have demonstrated that

organoaluminiums can be prepared by a Sn-Al exchange and alkenyl-, alkynyl-, allyl-, benzyl- and

arylaluminiums have been synthesized using alkylaluminium chlorides (Scheme 7).50

Organoboron reagents are relatively stable and most important, far less toxic than organomercury or

organotin compounds. Of course, because of the similarity between boron and aluminium, exchange

reactions between these two elements have been investigated.51 Thus, treatment of

triethylaluminium with triphenylboron at 140 °C leads rapidly to an exchange of ligands (Scheme

8).49e. equilibrium has to be shifted by removal of volatile reaction products, such as Et3B, to obtain

high exchange rates.54 Remarkably, this exchange has also been achieved using boracycles in which

case aluminium heterocycles are generated.

1.1.3 Hydrolalumination:

The addition of aluminium hydrogen bonds to unsaturated compounds such as alkenes and alkynes is

referred to as hydroalumination. The hydroalumination reaction was recognized by Ziegler51, 67

and has found application in one of the most important industrial syntheses of triorganoaluminiums:

the “Ziegler direct process” (Scheme 15).68 Two observations are combined in one process: First, Al

reacts with hydrogen in the presence of AlR3 leading to aluminium hydrides (increase) and secondly

the Al-H bond can react with alkenes (attachment). In summary, R3Al has been synthesized from Al,

H2 and alkene, the process therefore is referred to as “direct synthesis”.

The hydroalumination of alkynes and alkenes has developed to a general method for the preparation

of various alkenylaluminium and alkylaluminium reagents.70 Most recently it could be shown that

the use of Ni-salts catalyses the hydroalumination. Remarkably, by changing the Ni-catalyst,

regioselectivity of the hydrolalumination can be achieved.

2. Lead Organyls:

Organolead compounds are chemical compounds containing a chemical

bond between carbon and lead. The first organolead compound was hexaethyldilead (Pb2(C2H5)6),

first synthesized in 1858.

Synthesis:

Organolead compounds can be derived from Grignard reagents and lead chloride. For

example, methylmagnesium chloride reacts with lead chloride to tetramethyllead, a water-clear

liquid with boiling point 110 °C and density 1.995 g/cm³. Reaction of a lead(II) source

with sodium cyclopentadienide gives the lead metallocene, plumbocene.

Certain arene compounds react directly with lead tetraacetate to aryl lead compounds in

an electrophilic aromatic substitution. For instance anisole with lead tetraacetate forms 'p-

methoxyphenyllead triacetate in chloroform and dichloroacetic acid

Other compounds of lead are organolead halides of the type RnPbX(4-n), organolead sulfinates

(RnPb(OSOR)(4−n)) and organolead hydroxides (RnPb(OH)(4−n)). Typical reactions are:[4]

R4Pb + HCl → R3PbCl + RH

R4Pb + SO2 → R3PbO(SO)R

R3PbCl + 1/2Ag2O (aq) → R3PbOH + AgCl

R2PbCl2 + 2 OH− → R2Pb(OH)2 + 2 Cl−

R2Pb(OH)2 compounds are amphoteric. At pH lower than 8 they form R2Pb2+ ions and with pH

higher than 10, R2Pb(OH)3− ions.

Derived from the hydroxides are the plumboxanes:

2 R3PbOH + Na → (R3Pb)2O + NaOH + 1/2 H2

(R3Pb)2O + R'OH → 1/n (R3PbOR')n - n H2O

Organolead compounds form a variety of reactive intermediates such as lead free radicals:

Me3PbCl + Na (77 K) → Me3Pb.

Me3Pb-Pb-Me3 → [Me2Pb]

[Me2Pb] + (Me3Pb)2 → Me3Pb-Pb(Me)2-PbMe3

Me3Pb-Pb(Me)2-PbMe3 → Pb(0) + 2 Me4Pb

3. Tin organyls:

Organotin compounds or stannanes are chemical compounds based

on tin with hydrocarbon substituents. Organotin chemistry is part of the wider field of organometallic

chemistry. The first organotin compound was diethyltin diiodide ((C2H5)2SnI2), discovered by Edward

Frankland in 1849.

Preparation:

i) Organotin oxides and hydroxides are common products from the hydrolysis of organotin halides.

Unlike the corresponding derivatives of silicon and germanium, tin oxides and hydroxides often adopt

structures with penta- and even hexacoordinated tin centres, especially for the diorgano- and

monoorgano derivatives. The group Sn-O-Sn is called a stannoxane. Structurally simplest of the

oxides and hydroxides are the triorganotin derivatives.

R3SnOH ⇌ R3SnOSnR3 + H2O

ii) The reaction of a Grignard reagent with tin halides.

4EtMgBr + SnCl4 → Et4Sn + 4 MgClBr

Properties:

iii) The symmetrical tetraorganotin compounds, especially tetraalkyl derivatives, can then be

converted to various mixed chlorides by redistribution reactions.

3 R4Sn + SnCl4 → 4 R3SnCl

R4Sn + SnCl4 → 2 R2SnCl2

R4Sn + 3 SnCl4 → 4 RSnCl3

iv) The mixed organo-halo tin compounds can be converted to the mixed organic derivatives, as

illustrated by the synthesis of dibutyldivinyltin.

Bu2SnCl2 + 2 C2H3MgBr → Bu2Sn(C2H3)2 + 2 MgBrCl

References

1. Main Group Metals in Organic Synthesis Yamamoto, Hisashi / Oshima, Koichiro (eds.)

2004 ISBN 3-527-30508-4

2. Synthesis of Organometallic Compounds: A Practical Guide Sanshiro Komiya Ed. 1997

3. Caseri, Walter (2014). "Initial Organotin Chemistry". Journal of Organometallic

Chemistry. 751: 20–24. doi:10.1016/j.jorganchem.2013.08.009

School of Science and Humanities

Department of Chemistry

SCYA5201

Chemistry of Main Group Elements

Unit 5

LANTHANIDES & ACTINIDES

1. Actinides:

The two series of elements in the f block derive from the filling of the seven 4f and 5f orbitals,

respectively. This occupation of f orbitals from f1 to f14 corresponds to the elements cerium

(Ce) to lutetium (Lu) in Period 6 and from thorium (Th) to lawrencium (Lr) in Period 7. The

4f elements are collectively the lanthanoids (formerly and still commonly ‘the lanthanides’, but

this terminology is potentially confusing as the suffix -ide generally refers to an anionic form

of an element) and the 5f elements are the actinoids (still commonly the ‘actinides’). The

lanthanoids are sometimes referred to as the ‘rare earth elements’. That name, however, is

inappropriate because they are not particularly rare, except for promethium, which has no stable

isotope. A general lanthanoid is represented by the symbol Ln and an actinoid by An.

1.1CHARACTERISTICS OF THE LANTHANIDES

The lanthanides exhibit a number of features in their chemistry that differentiate them from

the d-block metals. The reactivity of the elements is greater than that of the transition metals,

akin to the Group II metals:

1. A very wide range of coordination numbers (generally 6–12, but numbers of 2, 3 or 4

are known).

2. Coordination geometries are determined by ligand steric factors rather than crystal field

effects.

3. They form labile ‘ionic’ complexes that undergo facile exchange of ligand.

4. The 4f orbitals in the Ln3+ ion do not participate directly in bonding, being well shielded

by the 5s2 and 5p6 orbitals. Their spectroscopic and magnetic properties are thus largely

uninfluenced by the ligand.

5. Small crystal-field splitting and very sharp electronic spectra in comparison with the d-

block metals.

6. They prefer anionic ligands with donor atoms of rather high electronegativity (e.g. O,

F).

7. They readily form hydrated complexes (on account of the high hydration energy of the

small Ln3+ ion) and this can cause uncertainty in assigning coordination numbers.

8. Insoluble hydroxides precipitate at neutral pH unless complexing agents are present.

9. The chemistry is largely that of one (3+) oxidation state (certainly in aqueous solution).

10. They do not form Ln=O or Ln≡N multiple bonds of the type known for many transition

metals and certain actinides.

Lanthanides: lanthanide series, abundance and natural isotopes - lanthanide contraction, similarity in

properties, occurrence, oxidation states - chemical properties of Ln(III) cations - magnetic properties, colour

and electronic spectra of lanthanide compounds - separation of lanthanides, solvent extraction, ion exchange

method.

Actinides: actinide series, abundance and natural isotopes, occurrence, preparation of actinides - oxidation

states - general properties, the later actinide elements, uranium-occurrence, metallurgy; chemical properties

of hydrides, oxides, and halides - complexes of lanthanides and actinides - term symbols

11. Unlike the transition metals, they do not form stable carbonyls and have (virtually) no

chemistry in the 0 oxidation state.

1.2 THE OCCURRENCE AND ABUNDANCE OF THE LANTHANIDES:

The lighter lanthanides are more abundant than the heavier ones; secondly, that the elements

with even atomic number are more abundant than those with odd atomic number. Overall,

cerium, the most abundant lanthanide on earth, has a similar crustal concentration to the lighter

Ni and Cu, whilst even Tm and Lu, the rarest lanthanides, are more abundant than Bi, Ag or

the platinum metals. The abundances are a consequence of how the elements were synthesized

by atomic fusion in the cores of stars with heavy elements only made in supernovae. Synthesis

of heavier nuclei requires higher temperature and pressures and so gets progressively harder as

the atomic number increases. The odd/even alternation (often referred to as the Oddo–Harkins

rule) is again general, and reflects the facts that elements with odd mass numbers have larger

nuclear capture cross sections and are more likely to take up another neutron, so elements with

odd atomic number (and hence odd mass number) are less common than those with even mass

number. Even-atomic-number nuclei are more stable when formed.

1.3 THE LANTHANIDE CONTRACTION:

As the series La–Lu is traversed, there is a decrease in both the atomic radii and in the radii of

the Ln3+ ions, more markedly at the start of the series. The 4f electrons are ‘inside’ the 5s and

5p electrons and are core-like in their behaviour, being shielded from the ligands, thus taking

no part in bonding, and having spectroscopic and magnetic properties largely independent of

environment. The 5s and 5p orbitals penetrate the 4f subshell and are not shielded. The

Lanthanides – Principles and Energetics increasing nuclear charge, and hence because of the

increasing effective nuclear charge they contract as the atomic number increases. Some part

(but only a small fraction) of this effect has also been ascribed to relativistic effects. The

lanthanide contraction is sometimes spoken of as if it were unique. It is not, at least in the way

that the term is usually used. Not only does a similar phenomenon take place with the actinides

(and here relativistic effects are much more responsible) but contractions are similarly noticed

on crossing the first and second long periods (Li–Ne; Na–Ar) not to mention the d-block

transition series. However, as will be seen, because of a combination of circumstances, the

lanthanides adopt primarily the (+3) oxidation state in their compounds, and therefore

demonstrate the steady and subtle changes in properties in a way that is not observed in other

blocks of elements. The lanthanide contraction has a knock-on effect in the elements in the 5d

transition series. It would naturally be expected that the 5d elements would show a similar

increase in size over the 4d transition elements to that which the 4d elements demonstrate over

the 3d metals. However, it transpires that the ‘lanthanide contraction’ cancels this out, almost

exactly, and this has pronounced effects on the chemistry, e.g., Pd resembling Pt rather than

Ni, Hf is extremely similar to Zr.

1.4 CAUSE OF LANTHANIDE CONTRACTION:

As we move along the lanthanide series from Ce to Lu, the addition of electrons takes place to

the 4f-orbitals, one at each step. The mutual shielding effect of f-electrons is very little, being

even smaller than that of d-electrons, due to the scattered or diffused shape of these orbitals.

However, the nuclear charge (i.e. atomic number) goes on increasing by one unit at each step

(i.e., each next element). Thus, the attraction between the nucleus and the outermost shell

electrons also goes on increasing gradually at each step. The 4f-electrons are not able to shield

effectively the attraction of the nucleus (i.e. inward pull) for the electrons in the outer most

shell as the atomic number of lanthanide elements increases. This results in the increased

inward pull of the outer most electrons by the nucleus, finally causing the reduction in the

atomic or ionic size of these elements. The sum of the successive reductions gives the total

lanthanide contraction. It may be concluded that the lanthanide contraction among the 4f-

sereies elements and their ions takes place due to the poor shielding effect of 4f-electrons and

gradual increase in the nuclear charge.

SIMILARITY IN PROPERTIES:

There is very small change in the radii of lanthanides and hence their chemical properties are

quite similar. This makes the separation of these elements using the usual physical and

chemical methods difficult. Consequently, new methods like ion exchange technique, solvent

extraction etc. have now been used for their separation which are based on slight difference in

the properties like hydration, complex ion formation, etc.

OXIDATION STATES:

The characteristic oxidation state of the lanthanide elements is ± 3. The universal preference

for this oxidation state together with the notable similarity in size led to great difficulties in the

separation of the elements prior to the development of chromatographic methods. Despite their

propensity to form stable +3 cations. the lanthanides do not closely resemble transition metals

such as chromium or cobalt. The free lanthanide metals are more reactive and in this respect

are more similar to the alkali or alkaline earth metals than to most of the transition metals. They

all react with water with evolution of hydrogen. The oxidation state (+III) is ionic and Ln+3

dominates the chemistry of these elements. The Ln2+ and Ln4+ ions that do occur are always

less stable than Ln3+. In the same way as for other elements the higher oxidation states occur

in the fluorides and oxides and the other halides. Particularly bromides and iodides. Oxidation

numbers (+II) and (+IV) do occur. Particularly when they lead to:

1. A noble gas configuration

2. A half-filled f shell

3. A completely filled f shell

The (+III) state is always the most common and most stable. The only (+IV) and (+II) states

which have any aqueous chemistry are Ce4+, Sm2+, Eu2+ and Yb2+. The lanthanide elements

resembles each other much more closely than do a horizontal row of the transition elements.

This is because the lanthanides effectively have only one stable oxidation state, (+III). Thus in

this series it is possible to compare the effects of small change in the size and nuclear charge

on the chemistry of these elements.

1.5 SPECTRAL PROPERTIES OF LANTHANIDES:

Most lanthanoid ions are only weakly coloured because their absorptions in the visible region

of the spectrum are commonly f-f transitions which are symmetry forbidden. The spectra of

their complexes generally show much narrower and more distinct absorption bands than those

of d-metal complexes. Both the narrowness of the spectral features and their insensitivity to

the nature of coordinated ligands indicate that the f orbitals have a smaller radial extension than

the filled 5s and 5p orbitals. The 5s and 5p orbitals are expected to shield the 4f electrons from

the ligands. f-orbitals are relatively deep inside the atom and overlap only weakly with ligand

orbitals. Hence, as a first approximation, their electronic states (and therefore electronic

spectra) can be discussed in the free-ion limit and the Russell-Saunders coupling scheme

remains a reasonable approximation despite the elements having high atomic numbers. the

visible spectra of lanthanoids usually consist of a large number of sharp, low intensity peaks

that are barely affected by changing the coordination environment of the central metal ion.

The spectra of the lanthanoid ions are normally characterized by the following properties:

1. Numerous absorptions due to the large number of microstates. Weak absorptions due

to lack of orbital mixing. Molar absorption coefficients(ε) are typically

2. 10 dm3 mol−1 cm−1 compared with d metals (close to 100 dm3 mol−1 cm−1).

3. Sharp absorptions due to the weak interaction of the f orbitals with the ligand vibrations.

4. Spectra that are to a large degree independent of the ligand type and coordination

number

1.6 MAGNETIC PROPERTIES:

Magnetism is a property associated with unpaired electrons. La3+ (f0), Ce4+ (f0 ), Yb2+ (f14)

and Lu3+(f 14) have no unpaired electron, hence they are diamagnetic. Other lanthanide ions

have unpaired electron in 4f orbital, so they are paramagnetic. Magnetic behaviour in case of

lanthanides arises due to the contribution of both spin moment as well as orbital magnetic

moment, unlike the d-block elements where magnetic moment corresponds to spin only values

[µ = n(n+2), n = no. of unpaired electrons]. Because in case of inner-transition f-block

elements, the 4f orbitals are quite shielded from the surroundings and do not interact strongly

with the ligands surrounding the metal ion. Consequently, the orbital motion of 4f orbitals are

not restricted (not quenched). Therefore, in case of lanthanides, the magnetic effect arising

from the orbital motion of the electron, as well as that arising from the electrons spinning on

its axis, contributes to the total magnetic moment. Therefore, magnetic moments of lanthanides

are calculated by taking into consideration spin as well as orbital contributions and a more

complex formula:

µ = 4S (S+1) + L (L+1) is used,

where L = orbital quantum no. and S = spin quantum no. In case of lanthanides, the spin

contribution (S) and orbital contribution (L) couple together to give a new quantum no. J, called

total angular momentum quantum no.

J = L - S, when the orbital is less than half full

J =L + S, when the orbital is less than half full

Magnetic moment values for lanthanides in Bohr Magneton (BM) is given by:

µ = g J (J + 1)

Where g is called Lande splitting factor and is given by:

1.7 SEPARATION OF LANTHANIDES BY ION-EXCHANGE METHOD:

This is the most rapid and most effective method for the isolation of individual lanthanide

elements from the mixture. An aqueous solution of the mixture of lanthanide ions (Ln3+aq) is

introduced into a column containing a synthetic cation exchange resin such as DOWAX-50

[abbreviated as HR (solid)]. The resin is the sulphonated polystyrene containing-SO3H as the

functional group. As the solution of mixture moves through the column, Ln3+aq ions replace

H +ions of the resin and get themselves fixed on it:

Ln3+aq + 3H-resin → Ln (resin)3 + 3H+ aq

The H+ aq ions are washed through the column. The Ln3+aq. ions are fixed at different positions

on the column. Since, Lu3+aq. is largest (Lu3+ anhydrous is smallest and is hydrated to the

maximum extent) and Ce3+aq. is the smallest, Lu3+aq. ion is attached to the column with

minimum firmness remaining at the bottom and Ce3+aq. ion with maximum firmness remaining

at the top of the resin column. In order to move these Ln3+aq. ions down the column and recover

them, a solution of anionic ligand such as citrate or 2-hydroxy butyrate is passed slowly through

the column (called elution). The anionic ligands form complexes with the lanthanides which

possess lower positive charge than the initial Ln3+aq ions. These ions are thus displaced from

the resin and moved to the surrounding solutions as eluant- Ln complexes. For example, if the

citrate solution (a mixture of citric acid and ammonium citrate) is used as the eluant, during

elution process, NH4 + ions are attached to the resins replacing Ln3+aq. ions which form Ln-

citrate complexes:

Ln (resin)3 + 3NH4 + → 3NH4- resin + Ln3+aq

Ln3+aq + citrate ions → Ln-citrate complex

As the citrate solution (buffer) runs down the column, the metal ions get attached alternately

with the resin and citrate ions (in solution) many times and travel gradually down the column

and finally pass out of the bottom of the column as the citrate complex. The Ln3+aq cations

with the largest size are, eluted first (heavier Ln3+aq ions) because they are held with minimum

firmness and lie at the bottom of the column. The lighter Ln3+aq ions with smaller size are held

at the top of the column (with maximum firmness) and are eluted at last. The process is repeated

several times by careful control of concentration of citrate buffer in actual practice.

1.8 LANTHANIDE SEPERATION BY SOLVENT EXTRACTION:

Solvent extraction has come to be used for the initial stage of the separation process, to give

material with up to 99.9% purity. In 1949, it was found that Ce4+ could readily be separated

from Ln3+ ions by extraction from a solution in nitric acid into tributyl phosphate [(BuO)3PO].

Subsequently the process was extended to separating the lanthanides, using a non-polar organic

solvent such as kerosene and an extractant such as (BuO)3PO or bis (2-ethylhexyl) phosphinic

acid [[C4H9CH(C2H5)]2P=O(OH)] to extract the lanthanides from aqueous nitrate solutions.

The heavier lanthanides form complexes which are more soluble in the aqueous layer. After

the two immiscible solvents have been agitated together and separated, the organic layer is

treated with acid and the lanthanide extracted. The solvent is recycled and the aqueous layer

put through further stages. For a lanthanide LnA distributed between two phases, a distribution

coefficient DA is defined:

DA = [LnA in organic phase] / [LnA in aqueous phase]

For two lanthanides LnA and LnB in a mixture being separated, a separation factor βAB can be

defined, where,

βAB = DA/DB

β is very close to unity for two adjacent lanthanides in the Periodic Table (obviously, the larger

β is, the better the separation). In practice this process is run using an automated continuous

counter-current circuit in which the organic solvent flows in the opposite direction to the

aqueous layer containing the lanthanides. An equilibrium is set up between the lanthanide ions

in the aqueous phase and the organic layer, with there tending to be a relative enhancement of

the concentration of the heavier lanthanides in the organic layer. Because the separation

between adjacent lanthanides in each exchange is relatively slight, over a thousand exchanges

are used. This method affords lanthanides of purity up to the 99.9% purity level and is thus

well suited to large-scale separation, the products being suited to ordinary chemical use.

However, for electronic or spectroscopic use (‘phosphor grade’) 99.999% purity is necessary,

and currently ion-exchange is used for final purification to these levels. The desired lanthanides

are precipitated as the oxalate or hydroxide and converted into the oxides (the standard starting

material for many syntheses) by thermal decomposition. Various other separation methods

have been described, one recent one involving the use of supercritical carbon dioxide at 40◦C

and 100 atm to convert the lanthanides into their carbonates whilst the quadrivalent metals (e.g.

Th and Ce) remain as their oxides.

2. ACTINIDES:

2.1 CHARACTERISTICS:

Early in the actinide series, electrons in the 6d orbitals are lower in energy than there is 5f

orbitals. This is clear from the ground-state electronic configurations of the atoms, which show

that the 6d orbitals are filled before 5f. The 5f orbitals are starting to be filled at protoactinium,

and with the exception of curium, the 6d orbitals are not occupied again. The 5f orbitals are not

shielded by the filled 6s and 6p subshells as the 4f orbitals of the lanthanides are (by the

corresponding 5s and 5p subshells). Moreover, the energy gap between 5f n 7s2 and 5f n−1 6d

7s2 configurations is less than for the corresponding lanthanides. Not only are the 5f orbitals

less ‘inner orbitals’ in the sense that the 4f orbitals are for the lanthanides, and thus more

perturbed in bonding, but the near-degeneracy of the 5f, 6d, and 7s electrons means that more

outer-shell electrons can be involved in compound formation (and a wider range of oxidation

states observed). Thus, the first four ionization energies of Th are 587, 1110, 1978, and 2780

kJ mol−1 compared with respective values for Zr of 660, 1267, 2218, and 3313 kJ mol−1. So,

for the earlier actinides, higher oxidation states are available and, as for the d block, several are

often available for each metal [relativistic effects may also be important. However, as the 5f

electrons do not shield each other from the nucleus effectively, the energies of the 5f orbitals

drop rapidly with increasing atomic number, so that the electronic structures of the later

actinides and their ions become more and more like those of the lanthanides, whose chemistry

they thus resemble.

OXIDATION STATES:

The important oxidation states exhibited by actinides are compiled below in the tabular form.

Some of them are stable but most of these oxidation states are unstable. It may be seen from

these oxidation states that the +2 state is shown by Th and Am only in the few compounds like

ThBr2, ThI2, ThS, etc. The +3 oxidation state is exhibited by all the elements and it becomes

more and more stable as the atomic number increases. The +4 oxidation state is shown by the

elements from Th to Bk, the +5 oxidation state by Th to Am, the +6 state by the elements from

U to Am and the +7 state is exhibited by only two elements, viz., Np and Pu. Np in the +7 state

acts as an oxidising agent. The principal cations given by actinide elements are M3+, M4+ and

oxo-cations such as MO2+ (oxidation state of M = + 5) and MO22+ (oxidation state of M = +6).

The examples of oxo-cations are UO2+, PuO2+, UO22+ and PuO22+ which are stable in acid and

aqueous solutions. Most of the M3+ ions are more or less stable in aqueous solution. Np3+ and

Pu3+ ions in solution are oxidized to Np4+ and Pu4+ by air. The latter ions are further oxidized

slowly to UO22+ and PuO22+ by air. The lighter elements up to Am show variable oxidation

states, the maximum being for Np, Pu and Am, but the heavier elements show oxidation state

of +2 or +3.

2.2 MAGNETIC AND SPECTRAL PROPERTIES:

It has already been mentioned that the paramagnetic nature of the substances is due to the

presence of unpaired electrons. The actinide elements like lanthanides show paramagnetism in

the elemental and ionic states. Tetravalent thorium (Th4+) and hexavalent uranium (U6+) ions

are diamagnetic due to the absence of unpaired electrons. Th4+ = U6+ = Rn (Z= 86) structure

(diamagnetic, paired electrons). Since, actinides constitute second f-series, it is natural to

expect similarities with lanthanides (the first f series) in their magnetic and spectroscopic

properties. But there are some differences between the lanthanides and actinides. Spin-orbit

coupling is strong (2000-4000cm-1) in the actinides as happens in the lanthanides but because

of the greater exposure of the 5f-electrons, crystal field splitting is now of comparable

magnitude and J is no longer such a good quantum number. It is also noted that 5f and 6d

subshells are sufficiently close in energy for the lighter actinides to make 6d levels accessible.

As a result, each actinide compound has to be considered individually. This must allow the

mixing of J levels obtained from Russel-Saunders coupling and population of thermally

available excited levels. Accordingly, the expression µ = g J(J+1) is less applicable than for the

lanthanides and magnetic moment values obtained at room temperature are usually lower and

are much more temperature dependent than those obtained for compounds of corresponding

lanthanides. The electronic spectra of actinide compounds arise from the following three types

of electronic transitions:

(a) f-f transitions: These are Laporte (orbitally) forbidden but the selection rule in relaxed

partially by the action of crystal field in distorting the symmetry of the metal ion. Because the

actinides show greater field, hence the bands are more intense. These bands are narrow and

more complex, are observed in the visible and UV regions and produce the colours in aqueous

solutions of simple actinide salts.

(b) 5f-6d transitions: These are Laporte and spin allowed transitions and give rise to much

more intense bands which are broader. They occur at lower energies and are normally confined

to the UV region hence do not affect the colours of the ions.

(c) Metal to ligand charge transfer: These transitions are also fully allowed and produce broad,

intense absorptions usually found in UV region, sometimes trailing in the visible region. They

produce the intense colours which are characteristic of the actinide complexes. The spectra of

actinide ions are sensitive to the crystal field effects and may change from one compound to

another. It is not possible to deduce the stereochemistry of actinide compounds due to

complexity of the spectra. Most of the actinide cations and salts are coloured due mainly to f-f

transitions. Those with f0, f7 and f14 configurations are colourless. The colours of some of the

compounds in different oxidation states are given below:

NpBr3 : green; NpI3 : brown; NpCl4 : red-brown; NpF6 : brown

PuF3 : purple; PuBr3 : green; PuF4 : brown; PuF6 : red brown

AmF3 : pink; AmI3 : yellow; AmF4 : dark tan.

The coordination chemistry of actinides is more concerned with aqueous solutions. Because of

the wider range of oxidation states available in actinides, their coordination chemistry is more

varied. Most of the actinide halides form complex compounds with alkali metal halides. For

example, ThCl4 with KCl forms complexes such as K[ThCl5] and K2[ThCl6], etc. ThCl4 and

ThBr4 also form complexes with pyridine, e.g. ThCl4.py Chelates are formed by the actinides

with multidentate organic reagents such as oxine, EDTA, acetyl acetone, etc. The actinides

with small size and high charge have the greatest tendency to form complexes. The degree of

complex formation for the various ions decreases in the order: M4+ > MO22+ > M3+ > MO2+.

The complexing power of different anions with the above cations is in the order:

Monovalent anions : F- > NO2 - > Cl- Bivalent anions : CO32- > C2O4

2- > SO42-.

2.3 SYNTHESIS:

Apart from thorium, protactinium, and uranium, the actinides are obtained by a bombardment

process of some kind. Thus for actinium (and also protactinium):

226Ra + 1 n → 227Ra + γ ; 227Ra → 227Ac + β− 230

Th + 1 n → 231Th + γ ; 231Th → 231Pa + β−

Similarly, for the elements directly beyond uranium:

238U + 1 n → 239U + γ ; 239U → 239Np + β−; 239Np → 239Pu + β−

and 238U + 1 n → 237U + 21 n; 237U → 237Np + β−

For obvious reasons, this route is not likely to be followed in the future; instead, heavy ion-

bombardment, using particles such as 11B, 12C, and 16O, will be used, though more recently

heavier ones like 48Ca and 56Fe have been utilized. This route is reliable but has the twin

drawbacks (from the point of view of yield) of requiring a suitable actinide target and also

being an atom-at-a-time route. Examples include:

244Cm + 4 He → 247Bk + 1 H

253Es + 4 He → 256Md + 1 n

248Cm + 18O → 259No + 3 1 n + 4 He

249Cf + 12C → 255No + 2 1 n + 4 He

248Cm + 15N → 260Lr + 3 1 n

249Cf + 11B → 256Lr + 4 1 n

HALIDES

The compounds known are summarized in Table 10.1. The only compound of an early actinide

in the +2 state is ThI2, a metallic conductor which is probably Th4+ (e−)2 (I−)2. Certain

heavier actinides form MX2 (Am, Cf, Es), which usually have the structure of the

corresponding EuX2 and are thus genuine M2+ compounds. All four trihalides exist for all the

actinides as far as Es, except for thorium and protactinium. Tetrafluorides exist for Th–Cm and

the other tetrahalides as far as NpX4 (and in the gas phase in the case of PuCl4). Pentahalides

are only known for Pa, U, and Np; whilst there are a few MF6 (M = U–Pu), uranium is the only

actinide to form a hexachloride. The known actinide halides are generally stable compounds;

most are soluble in (and hydrolysed by) water.

SIMILARITIES BETWEEN LANTHANIDES AND ACTINIDES

In the atoms of the elements of both the series, three outermost shells are partially filled

and remaining inner shells are completely filled but the additional or differentiating electron

enters (n-2) f-subshell.

The elements of both series exhibit +3 oxidation state which is prominent and

predominant state.

Like Lanthanide contraction found in the lanthanide elements, there occurs contraction

in size in the actinide elements called actinide contraction. Both the contractions are due to

poor shielding effect produced by f-electron with increasing nuclear charge.

The elements of both the series are quite reactive and are electropositive.

The electronic absorption bands of the elements of both the series are sharp and appear

like lines. These bands are produced due to f-f transitions within (n-2)f-subshell though such

transitions are orbital forbidden.

Most of the lanthanide and actinide cations are paramagnetic.

The nitrates, perchlorates and sulphates of trivalent lanthanide and actinide elements

are soluble while the hydroxides, fluorides and carbonates of these elements are insoluble.

The lanthanide and actinide elements show similarity in properties among their series

though the lanthanides are closer among them sieves in properties as compared to actinides.

DIFFERENCES BETWEEN LANTHANIDES AND ACTINIDES

Lanthanides Actinides

1. From left to right along 4f series, the

decrease of atomic size is nor regular with

increase in atomic number. But decrease of

ionic radius is regular

1. From left to right along 5f series, the

decrease of atomic size and ionic radius is

regular.

2. Along with +3 oxidation state, +2 and +4

oxidation state are also observed.

2. Along with +3 oxidation state, +4, +5, +6

and +7 oxidation state are also observed.

3. Lanthanide compounds are less basic. 3. Actinide compounds are more basic.

4. Ability to form complex compound is low. 4. Ability to form complex compound is high.

5. They does not form oxocation. 5. They form oxocation.

6. Except Pm, other lathanides are non-

radioactive.

6. All actinides are radioactive

References:

1. Textbook of Inorganic chemistry by J.D. Lee.